chemical element with symbol H and atomic number 1; lightest and most abundant substance in the universe

Hydrogen is a chemical element at the start of the periodic table. It has the symbol H and atomic number 1. It also has a standard atomic weight of 1.008. This makes it the lightest element in the periodic table. In standard conditions, hydrogen is a diatomic gas with the formula H
, or dihydrogen.[8] In this state, hydrogen is also called hydrogen gas or molecular hydrogen. Hydrogen has no color, smell, or taste.[9] Hydrogen is not toxic and is very combustible.[8]

Hydrogen, 1H
Purple glow in its plasma state
AppearanceColorless gas
Standard atomic weight Ar°(H)
Hydrogen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


Atomic number (Z)1
Groupgroup 1: hydrogen and alkali metals
Periodperiod 1
Block  s-block
Electron configuration1s1
Electrons per shell1
Physical properties
Phase at STPgas
Melting point(H
) 13.99 K ​(−259.16 °C, ​−434.49 °F)
Boiling point(H
) 20.271 K ​(−252.879 °C, ​−423.182 °F)
Density (at STP)0.08988 g/L
when liquid (at m.p.)0.07 g/cm3 (solid: 0.0763 g/cm3)[2]
when liquid (at b.p.)0.07099 g/cm3
Triple point13.8033 K, ​7.041 kPa
Critical point32.938 K, 1.2858 MPa
Heat of fusion(H
) 0.117 kJ/mol
Heat of vaporization(H
) 0.904 kJ/mol
Molar heat capacity(H
) 28.836 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 15 20
Atomic properties
Oxidation states−1, 0, +1 (an amphoteric oxide)
ElectronegativityPauling scale: 2.20
Ionization energies
  • 1st: 1312.0 kJ/mol
Covalent radius31±5 pm
Van der Waals radius120 pm
Color lines in a spectral range
Spectral lines of hydrogen
Other properties
Natural occurrenceprimordial
Crystal structurehexagonal
Hexagonal crystal structure for hydrogen
Speed of sound1310 m/s (gas, 27 °C)
Thermal conductivity0.1805 W/(m⋅K)
Magnetic orderingdiamagnetic[3]
Molar magnetic susceptibility−3.98×10−6 cm3/mol (298 K)[4]
CAS Number12385-13-6
1333-74-0 (H
DiscoveryHenry Cavendish[5][6] (1766)
Named byAntoine Lavoisier[7] (1783)
Isotopes of hydrogen
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
1H 99.9855% stable
2H 0.0145% stable
3H trace 12.32 y β 3He
 Category: Hydrogen
| references

Hydrogen is the most common chemical element in the universe. Hydrogen is almost 75% of all normal (baryonic) matter by mass.[10] Most stars are made of mostly hydrogen. Hydrogen stars are made of hydrogen in a plasma state. On Earth, hydrogen is seen in water and organic compounds. Hydrogen's most common isotope has one proton and no neutrons. This isotope also has one electron orbiting around it.

Hydrogen began to form a second after the Big Bang. These hydrogens did not have any neutrons or electrons. The first neutral hydrogen with an electron would not form until 380,000 years later during the recombination epoch, when the universe was cold enough for hydrogens to attract electrons.[11]

Hydrogen is usually nonmetallic and can form covalent bonds with most nonmetals. These covalent bonds can create molecules such as water and other organic substances. Hydrogen is the main part of acid–base reactions. These reactions exchange protons in soluble molecules. In ionic compounds, ions can either be anions or cations. Hydrogen anions are negatively charged and are called hydrides. Hydrogen cations are positively charged and are written as H+
. Cations are also called protons (symbol p), because they are only made of a proton and nothing else.

Hydrogen gas was first made artificially in the 1700s. Henry Cavendish identified hydrogen gas as a distinct substance between 1766 and 1781.

Most hydrogen production is from steam reforming natural gas. Hydrogen has many industrial uses. Hydrogen can be used to process fossil fuels, hydrocrack, and produce ammonia.



Hydrogen is grouped as a reactive nonmetal. This is different from the other elements found in the first group of the periodic table, which are called alkali metals. Only the solid form of hydrogen should behave like a metal, though.

When hydrogen is by itself, it will normally bind with itself to make dihydrogen (H2). Dihydrogen is very stable because of its high bond-dissociation energy of 435.7 kJ/mol.[12]

At normal temperature and pressure, hydrogen gas (H2) has no color, smell, or taste.[13] It is also not poisonous. This is because it is a nonmetal and burns very easily.[source?] Hydrogen gas at this state also has low density and is not corrosive.[13]



Molecular hydrogen is flammable and reacts with oxygen:

2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol)

At temperatures higher than 500 °C, hydrogen suddenly burns in air. This is called hydrogen autoignition temperature.[14]



While hydrogen gas in its natural form is not reactive, it does form compounds with many elements, especially halogens, which are very electronegative, meaning they want an electron very badly. Hydrogen also forms massive arrays with carbon atoms, forming hydrocarbons. The study of the properties of hydrocarbons is known as organic chemistry.

The H- anion (negatively charged atom) is named a hydride, though the word is not commonly used. An example of a hydride is lithium hydride (LiH), which is used as a "spark plug" in nuclear weapons.

Acids dissolved in water normally contain high levels of hydrogen ions, in other words, free protons. Their level is generally used to determine its pH, that is, the content of hydrogen ions in a volume. For example, hydrochloric acid, found in people's stomachs, can dissociate into a chloride anion and a free proton, and the property of the free proton is how it can digest food by corroding it.

Though uncommon on Earth, the H3+ cation is one of the most common ions in the universe.



Hydrogen has 7 known isotopes, two of which are stable (1H and 2H), which are commonly named protium and deuterium. The isotope 3H is known as tritium, has a half-life of 12.33 years, and is produced in small amounts by cosmic rays. The 4 isotopes left have half-lives on the scale of yoctoseconds.

Hydrogen in nature


In its natural form on Earth, hydrogen is generally a gas. Hydrogen is also one of the parts that make up a water molecule. Hydrogen is important because it is the fuel that powers the Sun and other stars. Hydrogen makes up about 74% of the complete universe.[15]

Natural hydrogen is normally made of two hydrogen atoms connected together. Scientists name these diatomic molecules. Hydrogen will have a chemical reaction when mixed with most other elements, though it has no color or smell.

Natural hydrogen is very uncommon in the Earth's atmosphere, because nearly all primordial hydrogen would have escaped into space because of its weight. In nature, it is generally in water. Hydrogen is also in all living things, as a part of the organic compounds that living things are made of. In addition, hydrogen atoms can join with carbon atoms to form hydrocarbons. Petroleum and other fossil fuels are made of these hydrocarbons and commonly used to make energy.

Some other facts about hydrogen:

History of Hydrogen


Hydrogen was first separated in 1671 by Robert Boyle.[18] In 1776, Henry Cavendish identified it as its own element and named it "inflammable air". He saw in 1781 that burning it made water.[19][5][6]

Hydrogen Spectrum Test

Antoine Lavoisier gave Hydrogen its name, from the Greek word for water, 'υδορ (read /HEEW-dor/) and gennen meaning to "produce"[20] as it forms water in a chemical reaction with oxygen.[21]

Uses of Hydrogen


The most common uses are in the petroleum industry and in making ammonia by the Haber process. Some is used in other places in the chemical industry. A little of it is used as fuel, for example in rockets for spacecraft. Most of the hydrogen that people use comes from a chemical reaction between natural gas and steam.

Nuclear fusion


Nuclear fusion is a very powerful source of energy. It depends on forcing atoms together to make helium and energy, as in a star like the Sun, or in a hydrogen bomb. This needs a large amount of energy to get started, and is not easy to do currently. A big advantage over nuclear fission, which is used in today's nuclear power stations, is that it makes less nuclear waste and does not use a poisonous and uncommon fuel like uranium. More than 600 million tons of hydrogen undergo fusion every second on the Sun.[22][23]

Using hydrogen


Hydrogen is mostly used in the petroleum industry, to change heavy petroleum parts into lighter, more useful ones. It is also used to make ammonia. Smaller amounts are burned as fuel. Most hydrogen is made by a reaction between natural gas and steam.

The electrolysis of water breaks water into hydrogen and oxygen, using electricity. Burning hydrogen joins with oxygen molecules to make steam (natural water vapor). A fuel cell joins hydrogen with an oxygen molecule, releasing an electron as electricity. For these reasons, many people believe hydrogen power will replace other synthetic fuels in the future.

Hydrogen can also be burned to make heat for steam turbines or internal combustion engines. Like other synthetic fuels, hydrogen can be made from natural fuels such as coal or natural gas, or from electricity, and therefore represents a valuable addition to the power grid; in the same role as natural gas. Such a grid and infrastructure with fuel cell vehicles is now planned by a number of countries, such as Japan, Korea and many European countries. This lets these countries buy less petroleum, which is an economic advantage. The other advantage is that, used in a fuel cell or burned in a combustion engine as in a hydrogen car, the engine does not make pollution. Only water, and a small amount of nitrogen oxides, forms.


  1. "Standard Atomic Weights: Hydrogen". CIAAW. 2009.
  2. Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN 978-0123526519.
  3. Lide, D. R., ed. (2005). "Magnetic susceptibility of the elements and inorganic compounds". CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. ISBN 978-0-8493-0486-6.
  4. Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 978-0-8493-0464-4.
  5. 5.0 5.1 "Hydrogen". Van Nostrand's Encyclopedia of Chemistry. Wylie-Interscience. 2005. pp. 797–799. ISBN 978-0-471-61525-5.
  6. 6.0 6.1 Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 978-0-19-850341-5.
  7. Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 978-0-19-508083-4.
  8. 8.0 8.1 PubChem. "Hydrogen". pubchem.ncbi.nlm.nih.gov. Retrieved 2024-05-30.
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  10. "Origin of the Elements". www2.lbl.gov. Retrieved 2024-05-30.
  11. "Early Universe - NASA Science". science.nasa.gov. Retrieved 2024-05-30.
  12. Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
  13. 13.0 13.1 "Hydrogen Fundamentals". hysafe/org. Retrieved 2024-05-30.
  14. Patnaik, P. (2007). A Comprehensive Guide to the Hazardous Properties of Chemical Substances. Wiley-Interscience. p. 402. ISBN 978-0-471-71458-3. Archived from the original on 26 January 2021. Retrieved 3 September 2020.
  15. Cain, Fraser, Universe Today (November 7, 2016). "When was the first light in the universe?". Phys.org. Archived from the original on 18 December 2016. Retrieved 29 November 2016.{{cite web}}: CS1 maint: multiple names: authors list (link)
  16. "EIA.doe.gov - What is Hydrogen?". Archived from the original on 2009-02-05. Retrieved 2008-12-22.
  17. "The magic of syngas". chemrec.se. 2012. Archived from the original on 20 April 2012. Retrieved 7 March 2012.
  18. Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
  19. Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
  20. Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 978-0-19-508083-4.
  21. Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 978-0-19-850341-5.
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